Binary compounds ionic and covalent
Ionic compounds contain a metal and a non-metal. Covalent compounds contain two non-metals. How do you determine whether an element is a metal? Simply refer to your periodic table. Metals are listed to the left of the thick stair-case line on the right hand side of the periodic table. Non-metals are grouped to the right of that line. Another feature of the periodic table indicates that metallic behavior increases as you travel down a group and decreases as you go from left to right across a row.
This will be important when we get to naming covalent binary compounds. All binary compounds are named with the more metallic element listed first and the more electronegative element listed second. Replace the ending of the electronegative element with -ide and you are pretty much on your way to naming your compound.
Then write the nonmetal element name second replacing the last three letters with -ide. Metals in the transition group often have multiple oxidation numbers. If your metal falls in the transition groups make sure to include Sn and Pb , then use Roman numerals to identify the oxidation number. Then write your non-metal element second with an — ide ending. Covalent binary compounds are composed of two nonmetals. Identify the location of each element on the periodic table.
The inference we draw from this is that the atom wants to obtain a filled shell, and this it achieves by forming bonds. This can be done by either addition of electrons or removal of electrons. The noble gas atom already has a filled shell and does not need to indulge in bonding to achieve it.
Elements on the left side of the table, metals, will lose electrons to form positive ions; elements on the right hand side of the table, non-metals, will gain electrons. In both cases, a filled shell will result. Of course, we must recognize that the atom is now charged because the electron and proton counts are not equal.
Electron loss creates positive ions, and electron gain creates negative ions. In an ionic compound, a positive ion and a negative ion come together and form an ionic bond through the strong electrostatic interaction between the ions of opposing charge. It is essential to be able to predict the charge on an ion in order to predict the composition of compounds formed containing it.
We can use the periodic table to assist us in this. The table shows the periodic table with the charges of the ions shown. Note, that in this version, the SI scheme of 1 — 18 is used rather than the older 1A — 8A. We find a very strong correspondence between group number using the older scheme and ion charge. A compound is always neutral, and so charges of the ions in the compound must balance out.
We always  know the charges on the ions from the periodic table. So the next stage is to determine the correct ratio of ions that will produce charge neutrality. Basically the total number of positive charges must equal the total number of negative charges. We have shown that the periodic table can be used to predict ionic charges.
However, there are some elements that are not susceptible to this approach. Some of the heavier A-type elements like tin and lead show two ionic charge possibilities: The transition metals also show a high degree of variable ionic charges: You are not expected to remember all of these different ions, but be able to predict a composition if given the ion, and write the composition with the correct notation.
The ionic bonding model works very well for metals and non-metals, but for compounds made exclusively from non-metals, which dominate chemistry in terms of numbers, it fails completely.
This is because non-metals form negative ions and never positive ions. It would also be impossible to describe the bond between the atoms in the diatomic elements like F 2 , O 2 and N 2 using the ionic model. A covalent H—H bond is the net result of attractive and repulsive electrostatic forces.
The nucleus—electron attractions blue arrows are greater than the nucleus—nucleus and electron—electron repulsions red arrows , resulting in a net attractive force that holds the atoms together to form an H 2 molecule. The sharing of electrons effectively increases the electron count around the atom. Alone, each fluorine atom has seven electrons in the outer shell. Sharing two electrons in a single covalent bond means that each atom now appears to have eight — it has satisfied its octet demand.
The same principle applies to describing bonds between unlike atoms, such as hydrogen and oxygen in water.